Rates of reaction — collision theory and factors

The rate of a chemical reaction is how quickly reactants are converted into products. Understanding rates is crucial for industrial chemistry (maximising efficiency) and everyday applications (cooking, medicine, corrosion).

Collision theory

For a reaction to occur, particles must:

1. Collide with each other.
2. Have sufficient energy to overcome the activation energy barrier (the minimum energy needed for the reaction to occur — the activation energy, Eₐ).

Any factor that increases the frequency or energy of successful collisions increases the rate of reaction.

Four factors affecting rate

1. Concentration (or pressure for gases): Higher concentration → more particles per unit volume → more frequent collisions → greater rate.

2. Temperature: Higher temperature → particles have MORE kinetic energy → move faster → more frequent AND more energetic collisions → more particles exceed the activation energy → much greater rate. A rule of thumb: raising temperature by 10°C approximately doubles the rate.

3. Surface area: For solid reactants, breaking them into smaller pieces increases the surface area in contact with the other reactant. More surface area → more particles exposed → more frequent collisions at the surface → greater rate. Example: powdered marble reacts faster with acid than marble chips.

4. Catalysts: A catalyst provides an alternative reaction pathway with a lower activation energy. More particles have enough energy to react → greater rate. The catalyst is not consumed in the reaction (it can be used repeatedly).

Measuring rate of reaction

Rate can be measured by:

• Volume of gas produced over time (using a gas syringe or measuring cylinder over water).
• Mass lost over time (if a gas is produced, on a balance).
• Change in colour or turbidity (for precipitation reactions — cross-on-paper/disappearing-X method).

Rate = amount of product formed (or reactant consumed) ÷ time taken.

Reaction profiles (energy diagrams)

A reaction profile shows the energy of reactants and products:

• Exothermic reaction: products are at a lower energy level than reactants; energy is released.
• Endothermic reaction: products are at a higher energy level than reactants; energy is absorbed.

The activation energy is the difference between the reactant energy level and the peak of the curve. A catalyst lowers the activation energy (shown as a lower peak in the diagram).

Common CCEA questions

• Describing and explaining the effect of a factor on rate using collision theory.
• Interpreting graphs of volume of gas vs time (steeper gradient = faster rate).
• Identifying the effect of a catalyst on a reaction profile diagram.